Periodic trends

Question 1

What does the atomic radius show and what is it’s vertical and horizontal trend in the table?

Answer 1

The atomic radius is used to compare the sizes of atoms. It is given in picometers (pm)
It is the distance between a nucleus and the valence electrons
Across a period the atomic radius decreases
Down a group the atomic radius increases

Question 2

What are the three ways of measuring the atomic radius and when are they applied?

Answer 2

Covalent radius: half the distance between two nuclei. The atoms must be bonded.
Technically the sum of two covalent radii should equal the covalent bond length between two atoms.
Metallic radius: Similar to covalent radius
Van der Waal’s radius: half of the distance between the closest approach of two non-bonded atoms (noble gases)

Question 3

Explain the theory behind the atomic radius and why it decreases across a period and increases down a group.

Answer 3

Across a period the number of valence electrons increase (the nuclear charge increases too). The valence electrons remain on the same shell which means the electron shielding remains constant.
With more valence electrons present, the attraction between them and the nucleus becomes stronger.
This means the electrons are pulled closer to the nucleus and the atomic radius decreases
Down a group, there are more shielding electrons which reduce the attraction of the valence electrons and the nucleus –> larger radius

Question 4

What does isoelectronic mean?

Answer 4

Ions of different elements with the same number of electrons (same electron configuration)
e.g. Na+, Mg2+, N3-, O2-

Question 5

What does the ionic radius show (related to anions and cations) and what is it’s vertical and horizontal trend in the table?

Answer 5

• The ionic radius decreases as the number of protons increase.
• Cations lose electrons in their valence shell, the number of protons remain the same which means the attraction increases
• The ionic radius decrease across a period and the cations have a smaller radius than their parent atoms.
• Anions have more electrons than protons. There is less of an attraction as the electrons increase the repulsion between nucleus and valence shell

Question 6

What is the effective nuclear charge and what depends on the strength of the electrostatic attraction?

Answer 6

• It is the attraction felt by the valence electron (the number of valence electrons)
• The effective nuclear charge increases across a period but remains the same down a group
• There is an electrostatic attraction between the nucleus and electrons in an atom.
• The strength of this attraction depends on the atomic radius and the no. of shielding electrons

Question 7

What is the ionization energy, it’s associated formula and what is it’s vertical and horizontal trend in the table? Explain the reasoning.

Answer 7

• The first ionization energy is the energy required to remove an electron from the attraction of the nucleus of a neutral atom to form a cation (+1)
• When explaining what it is include: nuclear charge, atomic radius and electron shielding
• Basically the measure of attraction between electrons and nucleus. (energy change positive)
  X (g) –> X (g) + e-
• Increases across a period
• Decreases down a group
• The more ionization energy there is, the more attraction there is between the nucleus and valence electrons
• Non-metals rather gain electrons and form anions as opposed to metals and therefore require more energy (higher ionization energy)

Question 8

What are the exceptions to the trend of increasing ionization energy across a group?

Answer 8

• There is a decrease in ionization energy from beryllium (s-block) to boron (p-block) (consecutive)
• The electron configuration of Boron ends with 2p^1 .
• The p-orbital is further away from the nucleus than the s-orbital therefore less energy is required to remove the electron in the p-orbital
• Boron is only partly filled (unstable can therefore be removed more easily) and beryllium has a full orbital (more stable)
• Also applies to magnesium to aluminum

Question 9

What are the other two exception to the ionization energy trend? Explain why.

Answer 9

• Decrease in ionization energy from nitrogen to oxygen. Also from phosphorus to sulfur.
• Nitrogen has an electron in each p orbital (more stable). Oxygen has 2 electrons in the first orbital. Less energy is required to remove the double occupied orbital because it is less stable.
• Therefore oxygen requires lower ionization energy, hence a decrease

Question 10

What is electronegativity and what is it’s vertical and horizontal trend in the table?

Answer 10

• The attraction of an atom for a bonding pair of electrons, basically the tendency of an atom to form a negative ion (gain electrons).
• Non-metals rather form anions (gain another electron) as opposed to metals.
• Non-metals have a higher electronegativity than metals.
• Increases across a period because there is a stronger attraction between the nucleus and bonding electrons
• Decrease down a group, less attraction
• The more attraction between the valence electrons and the nucleus, the higher the electronegativity

Question 11

What is the relationship between the electronegativity, effective nuclear charge and atomic radius?

Answer 11

• When the electronegativity increases, the effective nuclear charge increases and the atomic radius decreases
• More attraction means more electronegativity
  The difference in electronegativity between the atoms in a compound determines the type of bonding that occurs. When the difference is up to 1.7 —> covalent bonds. When the difference is over 1.8 units —> ionic bonds

Question 12

What is electron affinity and what is it’s vertical and horizontal trend in the table?

Answer 12

• The first electron affinity is the energy released when an electron is added to an atom to from a anion (1-)
• The amount of energy liberated when an atom acquires an electron from outside.
• First electron affinity (energy change negative)
  F (g) + e- –> F- (g) (exothermic)
• Second electron affinity:
  F- (g) + e- –> F2- (g) (endothermic)
• Non-metals (the halogens the most) have the highest electron affinities. When they gain electrons, they achieve a full valence shell which is stable.
• Increases across a period (more exothermic)
• Decreases down a group

Question 13

What are the metallic trends?

Answer 13

Large atomic radii, low electronegativity, low ionization energy, low electron affinity.
Have the tendency to lose electrons and form cations

Question 14

What are the non-metallic trends?

Answer 14

Small atomic radii, high electronegativity, high ionization energy, high electron affinity
Have the tendency to gain electrons to form anions

Question 15

What is the metallic trend across the periodic table?

Answer 15

Metallic character decreases across a period
Metallic character increases down a group
The metallic character is related to the ionization energy (tendency to form a cation)

Question 16

How is the melting point related to the type of bonding?

Answer 16

Across a period bonding changes from metallic to giant covalent to molecular covalent
When the strength of the bonds increase, so does the melting point
Molecular covalent compounds also have inter-molecular forces (weak, easily broken)

Question 17

What is the trend in bonding?

Answer 17

Ionic structures (metals) contain electrostatic forces
Molecular compounds (non-metals) contain inter-molecular forces 
Remember, the type of bonding across a period is determined by the difference in electronegativity between bonding atoms (covalent below 1.7)
Metallic atoms have a higher difference in electronegativity (above 1.8)

Question 18

What is the trend of oxides across a period and what is the definition of amphoteric?

Answer 18

Oxides change from basic to amphoteric to acidic
Amphoteric: a substance that can act as an acid and a base e.g. Al2O3
Period three oxides link

Question 19

What are the differences and similarities of the properties of group 1 metals (alkali metals)?

Answer 19

Physical properties: soft, first three elements low density
Melting and boiling point decrease down the group (generally low)
Reactivity increases down the group
Elements undergo vigorous reactions with water which produces metal hydroxides and H2 gas.
The solution then becomes alkaline (high pH)

Question 20

What are the differences and similarities of the properties of group 17 non-metals (halogens)?

Answer 20

Halogens are diatomic, and have separate physical states, molar mass increases
Chlorine: green gas
Bromine: red/brown liquid (gives off vapor)
Iodine: grey-solid (gives off purple vapor)
melting and boiling point decrease down the group

Question 21

Be able to construct equations with oxides and water reacting.

Answer 21

Oxides react with water to form as OH solution when basic. When acidic e.g. H2SO4 or HNO3
exception:
3NO2 + H2O –> 2HNO3 + NO

Question 22

Explain displacement reactions with halogens and what they have to do with oxidizing agents.

Answer 22

The more reactive halogen displaces the ions of the less reactive halogen from a solution
Ionic equations include negative ions on both sides.
A colour change occurs from displacement reactions. To enhance these colour changes, cyclohexane is added.
Halogens are oxidizing agents which become weaker down the group (when reactivity decreases too) Fluorine is strongest oxidizing agent.
Oxidizing agents, oxides (more positive) another substance but are reduced (more negative) themselves

Question 23

How do alkali metals react with halogens?

Answer 23

All the alkali metals react vigorously with halogens to produce salts e.g. NaCl
Both groups are the most reactive in the table